Process for recovering sodium carbonate from brines



2 Sheets-Sheet 1 4 A. c. HouGHToN Pnocss Fon nEcovERmG sopIUM cARBoNATE FROM Brumes Filed'April 9, 1956 Mm@ 19,y 1940.

`March 19, 1940, v v A. QHQUGHTON 2,193,817.

PROCESSVFOR RECOVERING SODIUM CARBONATE FROM BRINES Patented Mar.l19, v1940 treuil y Fsica .PROCESS FOR RECOVERING SODIUM f CARBONATE FROM BRNES Alexis (lrHoughton, Bartlett, Calif., assigner to.. FredericA. Dakin, Boston, Mass., as trustee Application April 9, 1936, Serial No. '73,498y

canas.` (o1. .2s-3s) lization through the monohydrate or any otherl This invention relates to an improved process of recovering commercially pure sodium carbonate from the brines of Owens Lake, California, and

other similar brines containing sodium carbonate,y

`5 sodium sulphate, and other constituents, by a crystallization method strictly according to the principlesv of physical chemistry. By commercially purev sodium carbonate is meant the soda ash of'commerce containing 99 per cent or more which seeks to avoid certain objections fully set I i'orth therein which are inherent in the carbonation process, such crystallization process making a purer product .at a cheaper ycos'tthan by carbonation.y In these inventions sodiuml carbonate deoahydrate containing a small amount of sodium sulphate decahydrate, is .-crystallizedby .refrigerating the brines under controlled conditions, and this produce is freedl from the smallv -amount of sodium sulphate by melting the decahydrate and evaporating off a portion ci the water to crystallize' out a pure sodium carbonate monohydrate free from sodium sulphate. The invention to be here described is much simpler than the inventions of the above named and Vother patents, as it crystallizes out a puresodium carbonate from the rines direct in one crystallization instead of two. Also, with the exception oi the drying operation, it operates throughout at temperatures not far removed from average room temperatures and readily obtainable with ordinary cooling water or 40 very slight warming, obviating the necessity of any artificial refrigeration or of boiling solutions, making it extremely economical in thermal energy requirements for either heating Orcooling.

Briefly the principle of this invention consists in crystallizing sodium carbonate directly from a brine saturated with sodium chloride and containing both sodium carbonate and a. certain amount oi sodium sulphate, under such conditions of concentration and temperature control that only purer sodium carbonate heptahydrate free from sodium sulphate will cr stallize out in a commercial yield of from 5i) to 70 per cent of the contained sodium carbonate of the brine, and then drying this sodium carbonate heptahydrate to the anhydrous state without further recrystalof NazCDs, equivalent to 58% NazO, on whichv basis the greater part of the soda ash manufac` form of sodium carbonate. v

It is weil known from various studies published in the chemical literature that When sodium carbonate is crystallized as the decahydrate from solutions containing both sodium carbonate and sodium sulphate, that mixed crystals or solid solutions oi sodium` sulphate decahydrate 'in sodium carbonate decahydrate are formed, the proportion of sodium sulphate in the decahydrate crystals depending on the ratio oi sulphate to 'carbonate in the` original solution from which the crystallization takes place. Both Na2COs.l0I-I2O and Na2SO4.10H2O crystallize inthe same system -the monoclinic, and under the microscope the crystal forms oi the two compounds are scarcely distinguishable from each other to vthe eye with-l out accurate measurement fof the angles. In other words the two compounds are isomorphous, and according to the laws governing isomorphism it is to be expected that they might form mixed crystals or solid solutions ci one in the other, as is the case in numberless other examples in the laboratory and in naturally occurring minerals` in which the molecules are similarly constituted and which crystallize in the same form. E. Mitscherlich, the classic authorityon the phenomena of isomorphism, makes the statement that while substances of dii-ferent crystalline form cannot combine in other than ilXeol proportions, substances oi the same crystalline form can crystallize together in all proportions. Further, it is stated by Blasdale A,in rdiscussing the phenomena of solid solutions that changes in the degree of hydration of a salt are invariably associated with prominent changes in the crystalline form it assumes as well as in its physical properties including itsability to dissolve other zsalts. There are no convincingv illustrations lof solid solutions whose component salts are hydrated but hydrated differently. In general unless the two salts are capable oi existing in the same degree oi hydration over a similar range of temperatures, and unless the two hydrates possess similar crystallographic properties, a continuous series oi solid solutionsl is not possible. Now it is well known that sodium carbonate heptahydrate crystallizes in a diiierent system to that of the decahydrates of sodium carbonate and sodium sulphate, namely in the rhombic system. While therefore there is no positive statement inthe chemical literature that sodium carbonate heptahydrate does not form mixed crystals with sodium sulphate decahydrate, .it is seen that according to well established principles governing the phenomena of isomorphism and solid solutions, the heptahydrate could not be expected to form such mixed crystals, and it could therefore be confldently predicted a priori that sodium carbonate heptahydrate would not in fact form such mixed crystals, and this indeed is found to be the case, namely, that sodium carbonate heptahydrate crystallizes out entirely free from sodium sulphate when crystallized from solutions containing both carbonate and sulphate. While it cannot be claimed therefore that this fact of sodium carbonate heptahydrate not forming solid solutions of sodium sulphate decahydrate is in itself a discovery or invention, as it is entirely deducible from generally accepted chemical principles, this present invention uses this fact and consists in determining and outlining the conditions for the largest possible formation of the heptahydrate from Owens Lake brines saturated with sodium chloride according to a simple method highly adaptable to economical commercial operation.

In the patent literature covering crystallization processes for recovering pure sodium carbonate from Owens Lake brines, it has been proposed to rst refrigerate the brine to separate out a crop in commercial yield of sodium carbonate decahydrate contaminated with some sodium sulphate in order to separate the sodium carbonate from the sodium chloride of the brine, and then recrystallize this sodium carbonate as the heptahydrate from an aqueous solution containing principally sodium carbonate with some sodium sulphate but practically free from sodium chloride, by means of adjusting the solution as regards its concentration and temperature so that crystallization takes place in the stable range of the heptahydrate. The temperature range within which the heptahydrate is stable in solutions free or nearly free from sodium chloride, is however very small, being from 33 to 30 C. Such a process requires a temperature control both during crystallization and ltering ofi1.5 C., which is a very difiicult and delicate one for commercial operation. The necessity for rst separating a commercial crop in the form of the decahydrate is stated to be the supposed fact that the direct crystallization of sodium carbonate heptahydrate from Owens Lake brines saturated with sodium chloride is impossible, as, it is stated, the eld of this crystalline form disappears entirely in the presence of this high concentration of sodium chloride. This supposition I have found to be entirely erroneous, as will be fully brought out in this specification. Not only is sodium carbonate heptahydrate stable in Owens Lake brines saturated with sodium chloride, but it is stable over a longer range of temperature and concentration than in straight sodium carbonate solutions in the absence of sodium chloride. For instance, the heptahydrate is stable in Owens Lake brines saturated with sodium chloride, from temperatures between 23.6 C. and about 17.5 C., or a range of 6 C. as compared with a range of only 3 C. in the carbonate solutions free from chloride. And what is even more important, and which is the discovery which makes this invention really practical and commercially economical as compared to the previous art, the range of metastable saturation above the transition point of the heptahydrate to the monohydrate, and below the transition point of the heptahydrate to the decahydrate, seems to be `enormously greater in saturated solutions of sodium chloride than in the absence of this salt. Full advantage of this discovery is taken in this invention. For

example, the concentration of such brines in sodium carbonate and saturated with sodium chloride at the transition point of the heptahydrate to the monohydrate at a temperature of about 23.6 C. is 16.3% NazCOs, but in the metastable range of the heptahydrate above this point the sodium carbonate concentration may be run up to as high as 21% NazCOa at a temperature of 31 C. At the other end of the stable heptahydrate eld at the transition point of the heptahydrate to the decahydrate at a concenration of around 12.8% NazCOa and a temperature of 17.5 C., the metastable range of the heptahydrate may be run down to as low as 9% NazCOa at a temperature of 10 C. This means that a brine saturated in sodium chloride and containing 21% NazCOs at a temperature of 31 C. may be cooled down to 10 C. with a NazCOa content in the mother liquor of 9%, with a total cooling range of 21 C. as compared to only 3 C. with the straight sodium carbonate solution free from sodium chloride, and with the recovery of a crop of '70% of the original sodium carbonate of the brine in the form of absolutely pure heptahydrate. It is this discovery of the remakable extension possible of the heptahydrate field in the metastable portions of the upper and lower part of the solubtility curve which is the chief feature of this invention, as it makes a process which is highly practical and commercial, and extremely simple to operate.

The above facts will be made clearer by reference to Figures 1 and 2, and an explanation of how they were derived.

Figure 1 is a curve of the solubility of the three hydrates of sodium carbonate in the simple system Na2CO3-H2O. It is drawn from data taken from a very accurate study of this system available in the literature. The dotted portion of the curves beyond the transition points was the extent to which the curves were followed into the metastable regions in this study.

In Figure 2 are curves of the solubility oi the three hydrates in Owens Lake brines saturated or substantially saturated with sodium chloride. By substantially saturated I mean saturated to within two or three per cent of sodium chloride.

In what follows, the terms decahydrate and hetpahydrate unless otherwise specied, relate to sodium carbonate decahydrate, NazCO3.10H2O, and sodium carbonate heptahydrate,

Two varieties of the heptahydrate exist, crystallizing in different forms and with different solubilities. These are known as the alpha and beta forms. The alpha form crystallizes in the rhombohedral system. It is more soluble than the beta variety, is quite unstable even in the solid crystalline form separated from the mother liquor, and has not been encountered in this work. The heptahydrate dealt with here is entirely the beta form, crystallizing in the rhombic system The term monohydrate refers to sodium carbonate monohydrate, Na2CO3.H2O.

Some explanation is necessary as to how the curves of Figure 2 were derived, especially the remarkable upper metastable portion of the heptahydrate curve, and this explanation will also serve to elucidate the principles of the invention.

In this study of the crystallization of sodium carbonate from Owens Lake brines, the peculiar fact has been constantly noted and repeatedly checked, that in attempting to obtain the highest concentration of sodium carbonate in such brines in the laboratory, whenv sodium carbonate monohydrate was used to saturate the brine at around 28 C., the concentration could never be made to exceed about 16% NazCOa. For `some vtime this' was` thought to be the highest possible concentration of sodium carbonate in such brines, especially as it was observed in the naturally occurring sub-surface lorinesin the crystal body of Owens Lake (which are always in contact with a large excess of sodium carbonate hydrates in the solid phase), that in the early fall when the temperature of such brine reaches y,its highest point (26-27 C.) .the concentration of sodium carbonate in solution was never more than about 16.5% NagCOg, the highest concentrationy of NazCOg in these brines ever observed being 17%. When however Owens Lake brines were saturated in the laboratory with a slight excess ol?v anhy-` drous sodium carbonate and the temperature not allowed to go over from 28 to 31 C., brines oontaining 19% to 21% NazCOn and completely saturated with sodium chloride were obtained with the greatest ease. Moreover, in making such saturations with anhydrous sodium carbonate, it was observed that the excess sodium carbonate not dissolved Ywas in very denite crystalline form, whereas the original anhydrous sodium carbonate added was amorphous. Examination under the microscope showed definitely that such crystals were those of the heptahydrate. The 21% NazCOa brine was therefore saturated with the heptahydrate at 31 C., but of course highly supersaturated with respect to monohydrate. The reason therefore such brines are not obtainednaturally in the lake is that they are in the metastable field of supersaturation with respect to inonohydrate, and in the crystal body of the lake there is an unlimited time element for nuclei and crystals of monohydrate to be formed when the temperature goes tooA far above the transition point concerned, and so if supersaturation of monohydrate should occur it is soon .broken and the solubility falls down to that of the monohydrate at 'the particular temperature. In the laboratory however these 19% to 2.1% sodium carbonate brines `were found to be very stable, and could be keptunchanged a practically unlimited length of time provided the temperature was kept constant at 28 to 31 C. and monohydrate seed crystals excluded. When cooled few degrees below 28 C. a plentiful crystallization of the heptahydrate occurred. When warmed a few degrees above`3l C., cloudiness at once resulted and a plentiiulprecipitation of the monohydrate soon followed. In the crystal body where there is no adequate stirring mechanism for equalization of temperature throughout the mass, local overheating would occur and thus start the precipitation of monohydrate, which would then proceed throughout the mass and equilibrium reached with the monohydrate at the particular temperature.

I do not wish. to limit myself by saying that higher saturations of sodium carbonate in a solution containing sodium sulphate and saturated with sodium chloride cannot be obtained at temperatures above 31 C. By very careful work it might be possible to obtain a solution containing more than 21% sodium' carbonate' and saturated with sodium chloride, but itis obvious that the 'farther the metastable part of the heptahydrate curve deviates from the solubility curve of the monohydrate, and the higher the temperature, the more unstable such a solution will be, and the more chance there will be of spontaneously forming monohydrate crystals andtl'iuszbrealdn the .supersaturation' in this constituent. For 'this' reason no attempt was made to prepare stronger solutions of sodium carbonate in such brines, as the increasing instability at the higher temperatures and concentrations involved would render them unsuitable for commercial work, where a Lake brine of approximately the same composition as given, on page 1, lines 21 to 33, of U. S. Patent #1,759,361 previously mentioned, was saturated at around 24 C. with an excess of anhydrous sodium carbonate. The heat .of hydration causes the temperature to rise to about 28 C. This is complete in about iteen minutes, after which the temperature israised to 31 C. and held there for thirty minutes, and then the excess soda filtered ofi. Such brine contains around.

21% NazCOa and around 12% NaCljbeing also saturated withthislatter salt. Such brine on cooling only one or two degrees starts to crystallize out heptahydrate, and the cooling was continued down todiiierent temperatures at intervals of 2` C., holding the solution at a constant temperature within `f0.1? C. at each desired temperature i'or half an hour with constant agitation, and then filtering 01T a small sample for analysis of the mother liquor with proper precautions for separating the mother liquor from the crystals at the` exact temperature desired by having the lter in the main mass of the liquid being held at constant temperature. Consistent results which checked each other very closely were easily obtained. For the lower temperatures it was necessary to removey some .of the crystals, aS otherwise-the mixture got too thick for 'good temperature equalization. The curve could not be followed much lower than 10 C., as at around this point there was a sharp evolution of heat and the mixture got almost solid due to complete conversion to the decahydrate. f

The reason for theready extension of the hep-tahydrate curve into the metastable region below the transition point of Athe heptahydrate to the decahydrate is believed to be as follows: Sodium carbonate decahydrate has a great tendency to form supersaturated solutions. It hasl been noted that in cooling Owens Lake brines i containing from10% to k13% NazCOg, that unless the solution is seeded with crystals of Na2CO3.lOH2O at around a temperature of 18 C., supercooling of from 5 to as muchr as 9 C. Without any crystallization is quite common. It is for this reason that in U. S. Patentitljfl we adopted the practice of always seeding the cooling solution with deoahydrate crystals to guard against the sudden crystallization of a large quantity of the decahydrate from a strongly supercooled'solution with the resulting formation of iine crystals diicult to lter and wash. With the stronger brines containing 15% to 21% NaaCOg, however, a cooling of only one or two degrees below the temperature at which the brine is saturated with NazCOa, always causes crystallisation of the heptahydrate to start without seeding of any kind. The heptahydrate, therefore, does not tend to supersaturate in nearly the same degree as the decahydrate. This is quite in accordance with the chemical principle` that the more highly hydrated a salt is, and consequently the vgreater the vapor pressure of H2O such salt has, the more tendency it has to form supersaturated solutions. Prominent examples of this are borax, NazBiOmlOHzO. and Glaubers salt, Na2SO4.10H2O, which are noted for their strong tendency to form supersaturated solutions. The vapor pressure of H2O of the heptahydrate being less than that of the decahydrate, it has not the same tendency towards supersaturation. If then a clear brine containing around 12.8% NazCOa at a temperature of 17.5 C., or at the transition point of the heptahydrate to the decahydrate as regards temperature and concentration, is cooled in the absence of seed crystals of either heptahydrate or decahydrate, a considerable amount of supersaturation with respect to decahydrate will occur. For several degrees below the transition temperature of 17.5 C. then such a solution could be supercooled with respect to the decahydrate, until at that particular temperature the composition of the solution came on to the metastable portion of the heptahydrate curve, and if the degree of supersaturation possible for the heptahydrate is small, heptahydrate would crystallize out provided that decahydrate crystals or nuclei had been excluded. Or under such conditions if heptahydrate crystals were already present in the solution at 17.5 C., but not decahydrate crystals, crystallization along the metastable path of the heptahydrate curve would continue. Eventually, however, as the cooling continues, owing to the growing divergence of the decahydrate and the metastable -portion of the heptahydrate curve, as shown in Figure 2, the limit of supersaturation with respect to decahydrate would be reached and exceeded, and the solution would enter into the labile eld of supersaturation of the decahydrate, where nuclei of Na2CO3-10I-I2O crystals would spontaneously form. Once decahydrate crystals are formed, the composition of the solution would come down to the decahydrate curve at the particular temperature, the decahydrate being less soluble than the heptahydrate. Also under such conditions on account of the greater solubility of the heptahydrate, all the heptahydrate crystals in the mixture will eventually dissolve and recrystallize out as the decahydrate. 'Ihese phenomena occur at about 10 C. It is for these reasons, and also for the reason that in previous refrigeration work on these brines the temperatures have for the most part been carried down below 10 C. and down to 5 C., that the heptahydrate phenomena have been entirely missed by Workers in this eld.

The curve for the decahydrate solubility was obtained by taking a 16% NazCOa brine, saturated of course with sodium chloride, cooling this down to about C. to crystallize out heptahydrate, and then cooling further to 5 C., when the crystals changed .over completely to the decahydrate. A sharp evolution of heat was noted when cooling below 10 C. This is caused by the fact that the heat of crystallization of the heptahydrate is less than that of the decahydrate. In converting the entire mass of already formed heptahydrate to decahydrate, therefore, a considerable amount .of heat is evolved. This mixture at 5 C. containing entirely decahydrate crystals and mother liquor, was now gradually warmed up, and samples ltered for analysis at intervals of 2.5 C. with the proper precautions of ltering at the exact temperature of the solution. When the solubility curve of the decahydrate so determined cuts the heptahydrate curve, or in other words when the solubility of the decahydrate is exactly the same as that of the heptahydrate, that is the transition point of the two salts in this particular brine. It was observed that at temperatures above this transition point, the excess of sodium carbonate decahydrate crystals in suspension was changed to the heptahydrate, and the metastable part of the decahydrate curve above the transition point could not be followed. The reason for this is the narrow field of metastable supersaturation of the heptahydrate, as previously explained. Also in brine above 12.8% NazCOs and 17.5 C. in temperature, if an excess of decahydrate crystals is added, and also enough solid sodium chloride to maintain saturation in this constituent, there is an almost immediate conversion of the whole mass of deca-hydrate to heptahydrate. This fact may be utilized in purifying sodium carbonate decahydrate contaminated with sodium sulphate by providing the proper conditions for this conversion to heptahydrate to take place, as will be seen in the example given later.

Ihe monohydrate solubility curve was determined by taking a brine containing 19% to 21% NazCOs and saturated with sodium chloride, prepared as described at 31 C., and warming slightly to 40 C., when there is a plentiful precipitation of the monohydrate. This mixture containing monohydrate in excess in the solid phase, was then cooled to lower temperatures at intervals of 2 C., and samples of mother liquor ltered 01T and analyzed.

Having fully described the principles underlying my invention, I will now describe and give examples of how it may be carried out in commercial operation. Before doing this, however, it will be necessary to briefly describe the present condition of the crystal body of Owens Lake.

Before the city of Los Angeles diverted the water of Owens River from Owens Lake, the lake had a depth of water or dilute brine of about 30 feet. The contained salts were then in complete solution, no crystallization of any salts having then occurred. In 1912 this dilute brine, according to the most trustworthy published analysis of the Geological Survey Laboratory in Washington, had the following composition:

Specific gravity 1.098 NazCOa per cent-- 4.29 NaCl do 3.96 KCI d0 .44 Na2SO4 do 1.61 Na2B407 d0 .27 NazSiOs do .05

It will be seen that sodium carbonate is the largest single constituent. Since the diversion of the main source of fresh water into this lake in 1917, the brine body has dried down to its present condition, which in the crystal body proper is that of a crystal mass or skeleton 8 to 10 feet deep at its deepest part and permeated with saturated brine. Very careful studies made of the crystal body in a section 8 feet deep of the composition of the crystal mass at dii'erent depths, have shown the following composition, expressed in water free salts.

NaCl NagCOg Na2SO4 Percent Percent Percent Upper 2% feet 80 l0 5 Next 1% feet 00 15 25 Next l foot 35 25 40 Lower 3 feet 5 00 5 At about 30 inches below the surface of the crystal body there is a hard salt layer a few inches thick. Above this layer the brines and salts are tember.

designated surface;" belowk this, sub-surface brines and salts. The chief characteristic of the surface part ofthe crystal body is the high percentage of sodium chloride and the large fluctuations of temperature, exposed as it is to the wide variations in atmospheric temperature throughout the year. However in the Owens Lake region, with its elevation of 3550 feet above sea level, there is more cold weather than warm throughout the year, the really lhot weather only lasting two or three months ln July, August, and September. As a consequence of this, and also from the fact that the sub-surface part is more or less protected from the temperature iiuctuations of the surface, the temperatureof the sub-surface part does not vary so widely. Froml prolonged records it is known that the temperature of the sub-surface brines varies from 14 C. at their lowest in March or April, to a peak of 27 C. n Sep- 'Ihere is no doubt therefore that the sodium carbonate in the sub-surface salts remains in the more highly hydrated forms of the decahydrate or heptahydrate throughout the greater part ofthe year, and only in the months of July, August, and September, is there likelihoodvof any conversion to the monohydrate. It must also be taken into account in this connection that the large mass of highly hydrated sodium carbonate and sodium sulphate, which persists from the colder weather, serves as a balance wheel or .equalizer of temperature changes, as in dissolving these hydrates large quantities of heat are absorbed, which tends to hold the temperature down. It is necessary to bring these facts out, as they play an important part in carrying out this invention.

Preferably the commercial plant would be located on the shore at the nearest point to the deepest part of the crystal body, in order to have the shortest possible pipe line for pumping a mixture of brine and salts. The deepest part of the crystal body is not in the center of the lake, as the bottom contour lines crowd toward the western shore, the sink of the lake being located only two miles from the western shore. By suitable mechanical means a mixture of brine and salts would be pumped through this relatively short pipe line, preferably from'a level four feet or greater from the surface of the lake in order to secure as high a percentage ci sodium carbonate salts as possible, although even the surface salts contain a suilicient quantity of sodium carbonate crystals for the purpose but more of them would have to be used. This mixture would be pumped to the plant, well stirred in a vessel provided with a suitable agitator, and supplied with both warming and cooling coils to hold the temperature at about 28 C. to 31 C. If there is a sufcient excess of sodium carbonate decahydrate or heptahydrate present in the salts, saturation up to 19% to 21% NazCO3 will occur in thirty minutes, or an even shorter time. There will always be a large enough excess of solid sodium chloride in the salts to maintain saturation in this constituent, as the brine in contact with the crystals as pumped is already saturated with sodium chloride. The contents of the vessel are then allowed to settle at the temperature of 28 C. to 31 C., and the clear brine drawn off, or if desirable the whole mass may be ltered over a rotary or other iilter. In hot weather should the salts pumped contain monohydrate, the brine will only saturate up to around 16% NazCOs. Several alternatives are possible in this rather unlikely contingency. No further attempt may be made to prepare a stronger brine, but such 16% NazCOc brine may be cooled down to say 12 C. with strict precautions against undercooling, so that there is no danger of going over to the decahydrate, and a recovery of 50% of the sodium carbonate originally in the brine obtained as pure heptahydrate. I have found that brines as low as 15% to 16% NazCOz; if *cooled carefully at the start will crystallize out the hep- 'tahydrate without the necessity of seeding with heptahydrate crystals. Or the mass of salts and brine may be agitated asuiicient length of time at a temperature below the transition-point of the monohydrate to heptahydrate and the monohydrate of the salts thus be converted to heptahydrate, when the temperature may be raised to 28 to 31 C. and the 19% to 21% NazCOs saturation obtained. Or finally saturation with the monohydrate containing salts may be obtained up to 16% NazCOa, the brine separated from the salts, and then a sufcient portion of the separated or fresh salts dried so that the monohydrate is converted to anhydrous `sodium carbonate, and the 16% NazCOs brine treated with a su'icient quantity of the dried salts to obtain the 19% to 21% saturation in sodium carbonate.

Such a prepared brine may have approximately tlle following composition:

By NaCl equivalent is meant the total chlorine of the brine calculated to sodium chloride. It is to be understood that the brine contains some 2% or 3% of KCl, which is thus included in the NaCl equivalent.

1,582 gallons of this prepared brine is cooled from a temperature of 31 C. to 16 C. with good agitation and temperature equalization throughout the liquid to avoid local undercooling, thereby y'crystallizing out 4,732 pounds of pure NazC'OsHI-IZO, containing 2,162 pounds of NazCOs, suspended in 13,255 pounds of` mother liquor of approximately the following composition.

Y Per cent NazCOs 12.0 NazBzO-r 3.2. NazSOr 4.6 NaCl equiv 16.4

The heptahydrate crystals are separated from the mother liquor by suitable mechanical means, preferably by centrifuging as permitting a more positive separation of mother liquor and a more eiective washing, and the crystals washed with. Water until as free as practicable from adhering mother liquor impurities. There will be obtained crystals containing 2,000 pounds of pure sodium carbonate. The heptahydrate crystals are then dried in a suitable dryer to produce 2,000 pounds f of soda ash having approximately the following VThe invention may also be applied, but with stable field of the heptahydrate. 35Y

a smaller recovery of sodiuinvcarbonate, to naturally occurring brines less strongly concentrated in sodium carbonate, say from 16% down to 13% NazCOz, or perhaps lower. For brnes under NazCOa it might be necessary to seed the cooling solution with crystals of the heptahydrate. Such crystals for seeding may be made with the greatest ease in any desired quantity by adding anhydrous sodium carbonate to a 12% solution of sodium chloride, or to any Owens Lake or similar brine sub-stantially saturated with sodium chloride. The sodium carbonate would be added under such conditions as would prevent the heat of hydration from carrying the temperature of the solution over 31 C. Under these conditions all the excess of NazCOa over that dissolved to increase the saturation of the mother liquor to 21% NazCOs would crystalline out as NazCOsJlI-IZO. Or a lower temperature may be used for the saturation down to 18 C. With a lower percentage of sodium carbonate remaining in the mother liquor. In this manner as large a crop of heptahydrate crystals as desired may be very easily prepared and filtered oi for use, or a suspension of the crystals in the mother liquor may be used for the purpose of seeding without filtering them off. After such seeding the brine could be cooled down to 12 C. and still crystallize out only the heptahydrate. but the danger of going over to the decahydrate would be greater unless very careful cooling with strict avoidance of local undercooling was practiced, as one would then be working in the meta- Vacuum cooling, by which the brine would be cooled to the desired temperature by boiling it under a high vacuum and the cooling accomplished by the latent heat or" evaporation of water, is ideally suited for this operation, as it automatically avoids any undercooled surfaces through which heat is being abstracted with a large temperature difference, as with vacuum cooling the temperature cannot go lower than the boiling point of the solution under the vacuum employed.

Any low grade salts containing a sucient quantity of sodium carbonate either as the decahydrate, heptahydrate, or preferably as anhydrous sodium carbonate, may be used for preparing a brine containing sufficient concentration of NazCOa to crystallize out only heptahydrate on cooling. Such low grade salts ideally suited for this purpose are trona deposits of Owens Lake. These are chemically the sesquicarbonate of soda, NazCOaNaI-ICOaZI-IzO, of which there are large deposits on the marginal portions of Owens Lake. In parts of the lake they are sufciently uncontaminated with sand or mud to be harvested cheaply and dried by a cheap source of heat, such as waste heat, to an anhydrous condition. Such a material may have the following composition:

Per cent NazCOs 64.4 Na2B204 1.4 Na2SO4 6.3 NaCl 15.3 Sand and foreign matter 12.6

Total 100.0

A lslight excess of the calcined trona over the amount necessary to bring the content of the brine to 21% Na2CO3 is added to any brine substantially saturated with sodium chloride at a temperature sufficiently lloW that lthe heat-of hydration' will not carry the final temperature over 31 C., and the temperature maintained at 28 to 31 C. for a short time, with agitation, the clear brine decanted or ltered 01T, and then cooled as previously directed to crystallize out the heptahydrate. It is obvious also that it is not necessary to use Owens Lake brine. Cheap sodium chloride could be used for preparing a brine containing slightly less than 12% NaCl, and the calcined trona added to this.

Also an impure soda ash containing sulphate and even large quantities of sodium chloride may be used as the saturating material.

Another low grade sodium carbonate containing material available at certain times of the year for this invention is the product known to operators on Owens Lake as winter soda. This is a crude decahydrate of sodium carbonate, which is crystallized out in large quantities when a surface or sub-surface brine containing over 9% NazCO3 is stored in a relatively shallow vat or pond, and subjected to atmospheric cooling in the winter time. Under such conditions more than 50% of the contained sodium carbonate of the brine may be crystallized out as a crude decahydrate, when the mother liquor may be pumped or drained out of the vat, and the crude decahydrate harvested and transported to the plant. The crude solid decahydrate crystals may be used as such as the saturating material, together with enough solid sodium chloride to maintain saturation in NaCl, or the crude decahydrate may be melted in the plant by warming up to its melting point, and the resulting liquor filtered or allowed to thoroughly settle in order to separate any monohydrate crystals that might have been formed in the heating and melting operation. This clear liquied crude decahydrate is now added to a surface or sub-surface brine in quantities suii'icient or in slight excess to bring the concentration of the resulting brine up to 21% NazCOs, while at the same time a sufficient amount of solid sodium chloride is added with agitation to maintain saturation of the brine in NaCl, while the temperature of the mixture is maintained at 28 to 31 C. The excess sodium carbonate crystallizing out and the excess of sodium chloride above that necessary for saturation, may be settled or filtered 01T, and the 19% to 21% NazCOz brine cooled to crystallize out a commercial crop of pure heptahydrate as previously described.

The invention may also be applied to another low grade sodium carbonate containing material. When Owens Lake brine or brine of similar composition is evaporated to recover potash, large quantities of waste salts are obtained containing a considerable amount of free sodium carbonate. By free sodium carbonate is meant that not combined with sodium sulphate as in the double salt Na2COs.2Na2SO4 known as burkeite. Such salts may have the following average composition:

Per cent NazCOal-Izo 7.3 Na2CO3.2Na2SO4 33.3 NaCl 48.0 Moisture 11.4

Total 100.0

In recovering 100 tons of potassium chloride from such brines, there may be obtained as muchas 50 tons of sodium carbonate in the form of monohydrate in such salts,'which may be recovered in a pure state by this invention. ToV 1000 parts of salts of the above composition would be .added 150 parts of Water and the mixture stirred at 31 C., which Would extract practically all of the sodium carbonate monohydrate in 390 parts of a solution of approximately the composition:

as in the example given in U. S. Patent #1,353,- 275. 1000 parts of fresh salts would be dried to convert the monohydrate to anhydrous sodium carbonate, and added to the 390 parts of the brine containing 16% NazCOa and agitated at a temperaturel not to exceed 28 to 31 C. There would thereby be formed 450 parts of a 19% to 21% NazCOs brine saturated with sodiumv chloride and W in sodium sulphate, which would be cooled to form only pure heptahydrate. The mother liquor from the heptahydrate filtrate could be used for re-saturation in sodium carbonate.

Also sodium carbonate decahydrate contaminated with a small amount of sodium sulphate, such as the product obtained in U. S. Patents #1,759,361 and #1,853,275, may be puried through the heptahydrate much more economically than previously proposed, by taking advantageof the fact that above the transition point of the decahydrate to heptahydrate insolutions saturated with sodium chloride, decahydrate crystals in suspension go over to the heptahydrate, and the sodium sulphate in solid solution in the decahydrate crystals is' thus liberated to the mother liquor. For instance a sufficient quantity of either a natural lake brine or an artificially prepared brine of just sodium chloride, is taken, so that the iinal mixture of solution and, suspended crystals Will not be too thick to handle and iilter easily. Such brine may contain 13% NazCOz or lower, and about 14% NaCl, being saturated with these constituents at about 18 C. To thisI brine is added an excess of decahydrate crystals of approximately the following composition:

Per cent Na2C'O.10I-I2,0 90.0 NazSOi. 10H20 4.2 Moisture and other constituents 5.8

Total 100.0

Also it will be necessary to add a suicient amount of solid sodium chloride to maintain substantial saturation in sodium chloride, but no-t to exceed this, as three molecules of Water will be given up to the mother liquor by the decahydrate and ten molecules of Water from the Glaubers salt when the NazCOsHHzO is formed, and also Water will be added by the moisture present, and all this Water must be substantially saturated With sodium chloride. There will be an absorption of heat due to the difference cf the heats of crystallization of the decahydrate and heptahydrate, so it will be necessary to Warm the solution up tov 18 C. cr'slightly over. This temperature is maintained with thorough stirring until all the decahydrate has gone into solution and recrystallized out as the heptahydrate, when all the s cdium carbonate in suspension Will be in the form of pure 'heptahydrate free fromsodium sulphate.

This may be filtered off, washed practically free from mother'liquor, and dried as usual to pure anhydrous sodium carbonate. The mother liquor Will still be saturated with sodium chloride, or nearly` so, if sufficient solid NaCl has been added, and will still contain only 13% NazCOs, and4 the recovery of the sodium carbonate of the decahydrate crystals as pure heptahydrate should be fully 90% Obviously also the method may be applied to a brine containing sodium carbonate, some sodium sulphate, but little or no sodium chloride, by saturating such brine with solid NaCl at a temperature around 31 C. If a high concentration of sodium carbonate is present, the excess over and above the 21% NazCOs remaining in the mother liquor Will crystallize out aspure crystallizing out the heptahydrate does not exf ceed 6% Na2SO4, at which point therewould be i danger of precipitating NazSOnlOHzOf as such with the heptahydrate. If the ratio of sodium sulphate to sodium carbonate in the original solution is high enough so that more than about 6% Na2SO4 would be formed in the final mother liquor, the sodium sulphate of the original brine may be reduced to a permissible -gure according to the principles of U. S. Patent #1,353,275, and then the above procedure applied to obtain a pure heptahydrate. f

I do not Wish to limit myself entirely to the above examples. Various combinations are possible in which the above principles may be applied to particular conditions or to particular materials using the principles outlined by those skilled in the art.

The advantages of this invention overv those previously proposed are many, and I will briefly enumerate some of them.

1. It is very simple and highly adapted to easy and economical commercial operation, involving as it does one saturation at a moderate temperature, only one filtering and Washing operation instead of two as in previous inventions, and a lnal comparatively low temperature drying operation. 'I'he-expense of extra filters for monohydrate is' thus eliminated, and the investment cost of a plant correspondingly lowered. Also the labor for operating the monohydrate filter .is

eliminated.

2. It is very economical in expenditure of thermal energy for either heating or cooling. The cooling range being within the scope of ordinary cooling Water for the greater part of the year,

it avoids the necessity of artificial refrigeration,y

such as ammonia compressor or other refrigerating machinery, as is necessary in other crystallizing processes on these brines. This makes a very considerable saving in investment cost of the plant, and also' of` labor charge for running such machinery. Also the heat oi? crystallization of the heptahydrate is considerably less than that of the decahydrate, being 192.5 B. t. u. per pound of sodium carbonate crystallized as the heptahydrate as compared with 273.3 B. t. u. per pound of sodium carbonate crystallized as' 'the decahydrate. Also there is no heat of crystallization'of sodium sulphate decahydrate such as occurs to a small extent in the decahydrate process.

3. It avoids the use of evaporators for forming the monohydrate as in previous crystallization processes, thus saving in heat consumption for heating and evaporating the sodium carbonate solution. The heptahydrate may be dried directly iirst to monohydrate, and finally to anhydrous sodium carbonate, in the same dryer. The heptahydrate in crystal form readily parts with its water of crystallization down to the monohydrate. In the dryer a considerable portion of the heat supplied for converting monohydrate to anhydrous carbonate in the heated end of the dryer would be carried back to the part Where the heptahydrate would be fed, and six molecules of Water of the heptahydrate driven off by this mild heat, and the monohydrate progressed forward to the heated end of the dryer to be converted to anhydrous sodium carbonate. Thus the heptahydrate would be dried practically by Waste heat to the monohydrate instead cf consuming steam for heating a solution and evaporating oilD water. Another considerable saving in thermal energy is that in the decahydrate process the decahydrate must be melted before being fed to the evaporator, Theoretically the melting absorbs exactly the same amount of heat as was given off in its crystallization. The thermal energy therefore required to take care of the heat of crystallization in forming the decahydrate from a solution is exactly doubled by the necessity of having to melt it again. This double absorption of thermal energy is avoided in the heptahydrate process, as the heptahydrate is not melted, but fed directly to the dryer, Where the water of crystallization is driven off preferably while the sodium carbonate is all in the solid state. The elimination of evaporate-rs or making provision for melting the decahydrate also reduces the investment cost of the plant and labor cost for operating same.

4. There is no silica problem in this invention. In previous crystallization processes for Owens Lake brine, the crystallization of the decahydrate carried the alkalinity of the mother liquor below the point where the silica was held in solution, and sufhcient quantities of silica separated out to seriously contaminate the nal product. This was prevented in U. S. Patent #1,759,361 by maintaining the alkalinity of the mother liquor by the addition of caustic soda. In the heptahydrate process the alkalinity of the mother liquor, owing to its higher sodium carbonate content, does not go below the point Where silica starts to separate out. The silica problem therefore does not appear in this process.

5. Finally, a soda ash of light apparent density is produced directly by this process. When monohydrate is crystallized from a solution and dried to anhydrous sodium carbonate, it results in a so-called dense soda ash, of an apparent density of about 60 pounds to the cubic foot. While this is advantageous for certain uses, such as glass making, a certain proportion of the soda ash trade requires a. lighter ash of density around 30 pounds to the cubic foot. In the ammonia soda process the product of the calcination of the sodium bicarbonate naturally results in a so-called light ash, and the dense ash requirement of the trade is made from light ash by either incipient or complete fusion of the sodium carbonate, and grinding the product, or by adding water to the light ash in the approximate proportions to form the monohydrate, and drying this monohydrate to a dense ash. The problem, however, of making a true light ash from a dense ash, which appears in all crystallization processes to the present from this or similar brines, has so far not been satisfactorily solved. The problem is solved in this invention if certain conditions are observed. The apparent density of soda ash is a function largely of particle size. Monohydrate crystallized from solutions has a fairly large, compact crystal, with the result that when its water is driven off, the particles of soda ash are still sufficiently large and compact to give a true dense ash. If nely ground to pass a 400 mesh screen, dense ash can be made into light ash. The expense of this additional operation, however, is to be avoided if possible. When, however, either the heptahydrate or the decahydrate is allowed to lose its water of crystallization under conditions that no actual melting of the hydrate takes place, or in other words when it undergoes a true efflorescence, the resulting crystals of monohydrate are extremely minute, passing a 400 mesh screen. When such monohydrate is finally dried to the anhydrous form, the resulting soda ash is a true light ash. In this invention full advantage may be taken of this fact without any extra operation or expense. All that suffices is to see that the heptahydrate is first exposed to very gentle temperatures insuicient to melt same or to cause fritting, partial melting, or any solution in water, until a large proportion of the water is driven oi. As the vapor tension of Water of the heptahydrate is greater than the normal vapor tension of water in the atmosphere, especially in the desert region of Owens Lake, where the average humidity is very low, a current of relatively dry warm air o1' gases would effect this operation without any actual melting of the heptahydrate. In this manner a true light ash of as low a density as 25 pounds to the cubic foot can be obtained. This is not possible in the invention of U. S. Patent #1,853,275 without extra expense, as a1- though the decahydrate eilloresces as easily or even more readily than the heptahydrate and would thus produce a light ash if dried in this manner, the decahydrate is not pure sodium carbonate, and if dried directly without further purification would not produce a 58% Na20 commercially pure soda ash.

Dense ash can be made from the light ash in the same manner as in the ammonia soda industry. Or if desired a dense ash may be made directly by melting up the heptahydrate or dissolving it in a small amount of Water, and then evaporating this sodium carbonate solution to crystallize out the monohydrate, liltering and washing this monohydrate, and drying same directly to produce a dense ash.

What I claim is:

1. The method of crystallizing pure sodium carbonate heptahydrate consisting in the preparation of a solution by dissolving to saturation a crude, solid sodium carbonate in a brine saturated wth respect to sodium chloride and containing less than ve per cent. sodium sulphate at a temperature between 28 C. and 31 C., cooling said solution without any undercooling to a temperature not lower than 10 C. to crystallize sodium carbonate heptahydrate from said solution and then separating out the crystalline material so obtained.

2. The method of crystallizing pure sodium carbonate heptahydrate consisting in the preparation of a solution by dissolving to saturation a crude, solid sodium carbonate in a brine saturated with respect to sodium chloride and containing less than ve per cent. sodium sulphate carbonate heptahydrate from a sodium carbonate brine containing sodium chloride and also sodium sulphate in an amount such that the concentration of sodium sulphate in the final heptahydrate mother liquor will be less than 6%, which consists in agitating the brine having substantially the composition of an Owens Lake brine with a mixture of salts containing an excess of the solid phases of sodium chloride and any form of sodium carbonate except sodium carbonate monohydrate, while warming the mixture to a temperature not to exceed 31 centigrade,y whereby the prepared brine contains from 161/% to 21% of sodium carbonate and is supersaturated with respect to sodium carbonate monohydrate, removing undissolved matter from the brine, and cooling the Aclear brine from said solution temperature to a temperature above 10 centigrade under conditions of agitation and no local undercooling, thereby crystallizing out of the brine' from 50% to 70% of its sodium carbonate content in the form of pure heptahydrate crystals free from crystallized sodium sulphate and` all other crystallized impurities.

fl. The method of crystallizing pure sodium carbonate heptahydrate from brines substantially saturated with sodium chloride and containing from 13% to ll/2% of sodium carbonate and an amountofv sodium sulphate such that the concentration of this constituent in the final mother liquor shall be'less than 6%, which consists in first coolingv such brines to a temperature between 18 and 24 centigrade, then seeding such brines with crystals of sodium carbonate heptahydrate and continuing the cooling to a temperature about 10 centigrade under conditions agitation and temperature equalization which avoid undercoollng lanywhere throughout the` mixture in order to prevent the formation of sodium carbonate decahydrate and to crystallize out only pure sodium carbonate heptahydrate free from crystallized sodium sulphate and any other crystallized impurities'in a yield of from 40% to 50% o1' the sodium carbonate in the brine.

5. A method of obtaining a high yield of pure sodium carbonate from crude sodium carbonate decahydrate crystals, which comprises melting such crystals, separating undissolved matter from the liquor thus formed, adding a suicient quantity yofthe clear liquor to an Owens Lake brine which is maintained at a temperature of from 28 to 31 C. to bring the sodium carbonate concentration in the brine to approximately 21%, the mixed brine also being maintained saturated with sodium chloride, separating undissolved matter from the mixed brine, and cooling the resultant brine from the solution temperature to a temperature above 10 C., vwith careful avoidance of local under-cooling, to crystallize out of the solution pure crystals of sodium carbonate heptahydrate.

ALEXIS C. HOUGHTON. 

